DAT General Chemistry: States of Matter and the Six Phase Changes

DAT High-Yield Concept

Thermodynamic Sign Conventions: Endothermic transitions ( $\Delta H > 0$ ) break intermolecular forces (fusion, vaporization, sublimation), whereas exothermic transitions ( $\Delta H < 0$ ) form intermolecular bonds (freezing, condensation, deposition).
Phase Transitions Skip Liquids: Sublimation transitions directly from solid to gas ( $\ce{CO2(s) -> CO2(g)}$ ); deposition transitions directly from gas to solid ( $\ce{CO2(g) -> CO2(s)}$ ).
Hess's Law for Enthalpy: Sublimation enthalpy equals fusion enthalpy plus vaporization enthalpy: $\Delta H_{\text{sub}} = \Delta H_{\text{fus}} + \Delta H_{\text{vap}}$ .
Physical Properties: Gases have indefinite shape/volume and high compressibility; liquids have definite volume but indefinite shape; solids have fixed shape, volume, and minimal compressibility.

Characteristics of Solids, Liquids, and Gases

Matter exists in three classical states determined by the competition between particle kinetic energy and attractive intermolecular forces (IMFs):

Solids: Intermolecular forces dominant over kinetic energy. Particles occupy fixed lattice points and vibrate in place, yielding definite shape, definite volume, and near-zero compressibility.
Liquids: Kinetic energy and intermolecular forces are balanced. Particles remain in close contact but move past one another, producing a definite volume with an indefinite shape that conforms to its container.
Gases: Kinetic energy far exceeds attractive forces. Particles undergo rapid, continuous translation with minimal interactions, resulting in indefinite shape, indefinite volume, and high compressibility.

Property	Solid	Liquid	Gas
Shape / Volume	Definite / Definite	Indefinite / Definite	Indefinite / Indefinite
IMF Strength	Very Strong	Moderate	Weak / Negligible
Particle Spacing	Tightly packed	Close together	Far apart
Compressibility	Negligible	Very low	High

Key Distinction: Temperature controls average kinetic energy (

E_k \propto T

). Heating increases kinetic energy to overcome IMFs, favoring gas phase; cooling permits IMFs to pull particles into condensed phases (liquid or solid).

The Six Phase Changes

A phase transition is a physical process in which a substance converts from one state of matter to another without altering its chemical identity:

Fusion (Melting): Transition from solid to liquid, e.g., $\ce{H2O(s) -> H2O(l)}$ .
Freezing (Solidification): Transition from liquid to solid, e.g., $\ce{H2O(l) -> H2O(s)}$ .
Vaporization: Transition from liquid to gas, e.g., $\ce{H2O(l) -> H2O(g)}$ . Includes surface evaporation and bulk boiling.
Condensation: Transition from gas to liquid, e.g., $\ce{H2O(g) -> H2O(l)}$ .
Sublimation: Direct transition from solid to gas, e.g., $\ce{CO2(s) -> CO2(g)}$ .
Deposition: Direct transition from gas to solid, e.g., $\ce{CO2(g) -> CO2(s)}$ .

The six phase changes, and which direction absorbs energy — Six changes connect the three states, and they pair up as opposites: melting with freezing, vaporization with condensation, sublimation with deposition. Going toward the more spread-out state always requires energy to be put in, because the attractions holding the particles together must be overcome; the reverse direction releases that energy. Sublimation and deposition skip the liquid entirely, which is what solid carbon dioxide does at ordinary atmospheric pressure: it sublimes straight to gas without ever melting.

Energetics: Endothermic vs. Exothermic Transitions

Phase transitions involve heat transfer between system and surroundings:

Endothermic ( $\Delta H > 0, \Delta S > 0$ ): Heat energy is absorbed to disrupt IMFs and increase particle disorder. Includes fusion ( $\Delta H_{\text{fus}}$ ), vaporization ( $\Delta H_{\text{vap}}$ ), and sublimation ( $\Delta H_{\text{sub}}$ ).
Exothermic ( $\Delta H < 0, \Delta S < 0$ ): Heat energy is released as IMFs form and particle order increases. Includes freezing ( $\Delta H_{\text{freeze}} = -\Delta H_{\text{fus}}$ ), condensation ( $\Delta H_{\text{cond}} = -\Delta H_{\text{vap}}$ ), and deposition ( $\Delta H_{\text{dep}} = -\Delta H_{\text{sub}}$ ).

Because enthalpy is a state function, Hess's Law dictates that direct sublimation enthalpy equals the sum of fusion and vaporization enthalpies:

\Delta H_{\text{sub}} = \Delta H_{\text{fus}} + \Delta H_{\text{vap}}

Memory Aids:
• END-O (Heat In,

\Delta H > 0

): Fusion, Vaporization, S**ublimation ("F**ind **V**aluable **S**unlight").
• EX-O (Heat Out, $\Delta H < 0$ ): Freezing, C**ondensation, D**eposition ("F**reeze **C**old **D**rops").

Worked Example: Thermochemical Phase Transition Calculation

Problem: Calculate the molar enthalpy of sublimation ( $\Delta H_{\text{sub}}$ ) for water given $\Delta H_{\text{fus}} = \pu{6.01 kJ/mol}$ and $\Delta H_{\text{vap}} = \pu{40.67 kJ/mol}$ . Classify the process $\ce{H2O(s) -> H2O(g)}$ as endothermic or exothermic.

Solution:

Write balanced chemical equations for fusion and vaporization: $\ce{H2O(s) -> H2O(l)} \quad \Delta H_{\text{fus}} = \pu{+6.01 kJ/mol}$ $\ce{H2O(l) -> H2O(g)} \quad \Delta H_{\text{vap}} = \pu{+40.67 kJ/mol}$

Add the equations, canceling liquid water $\ce{H2O(l)}$ : $\ce{H2O(s) -> H2O(g)}$

Calculate $\Delta H_{\text{sub}}$ : $\Delta H_{\text{sub}} = \Delta H_{\text{fus}} + \Delta H_{\text{vap}} = \pu{+6.01 kJ/mol} + \pu{+40.67 kJ/mol} = \pu{+46.68 kJ/mol}$

Conclusion: Since $\Delta H_{\text{sub}} = \pu{+46.68 kJ/mol} > 0$ , sublimation absorbs heat and is endothermic.

Common DAT Exam Trap: Students often mistake system vs. surroundings during freezing ( $\ce{H2O(l) -> H2O(s)}$ ). Freezing releases energy from system to surroundings ( $\Delta H_{\text{sys}} < 0$ ). Freezing water warms its immediate surroundings; it does not "absorb cold."

Critical Thinking Questions

Why does dry ice ( $\ce{CO2(s)}$ ) sublime directly into gas at room pressure ( $\pu{1 atm}$ ), whereas ice ( $\ce{H2O(s)}$ ) melts into liquid?

When perspiration ( $\ce{H2O(l)}$ ) evaporates from human skin, why is this process endothermic and how does energy transfer lower body temperature?

How does increasing intermolecular force strength affect a liquid's molar enthalpy of vaporization ( $\Delta H_{\text{vap}}$ ) and normal boiling point?

References

Sources returned by Google Search grounding during lesson verification.

11.4: Phase Changes - Chemistry LibreTexts — chem.libretexts.org
Endothermic and exothermic reactions - Khan Academy — khanacademy.org

States of Matter and the Six Phase Changes

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